Exp. Equil. Handout - Lab Report PDF

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2/23/18 Experiment: Equilibrium and Le Chatelier’s Principle

Purpose: The purpose of this experiment is to see how two equilibrium systems shift when a stress such as concentration and temperature changes occur. Theory/Principles: Equilibrium systems are when the rate of a forward and reverse reaction are equal for a reversible reaction. According to Henri Le Chatelier when a stress is applied to a system at equilibrium if will shift to minimize the effect of the applied stress. An example would be when increasing the concentration of a reactant it will cause the decrease of concentration for reactants while there will be an increase of concentration for products (a shift to the right). Another example would be when increasing temperature for a reaction mixture at equilibrium it will result in the reaction shifting to the right towards the products making it endothermic. Whereas, an exothermic reaction will shift to the left. In this experiment the first reaction we studied is CrO^2- 4 (aq) + 2H+ (aq) ⇌ Cr2O^2- 7 (aq) + H2O (l), we are told that CrO2- 4 solution has a more intense yellow color and when dichromate ion are present it will turn into an orange color. This procedure allows us to see how the concentrations of the chromate and dichromate ion change. We will also be adding Ba2+ ions to the solution for it to from a stable precipitate with one of the anions causing the removal of that ion to shift the equilibrium to minimize the loss. The second reaction we will study is [CoCl4]2- (alc) + 6H2O (l) ⇌ [Co(H2O)6]2+ (alc) + 4Cl(alc), in this reaction stresses are going to be placed where equilibrium systems can be monitored by observing color changes and measuring their absorbance’s. The observations we write down occurring from the stresses being placed on each system will be used to help us determine whether the reaction in each system is endothermic or exothermic. Experimental Procedures:

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Handout given to us by lecture professor / at the end of class syllabus

Anderson, V. K.; Mathias, E. V. Equilibrium and Le Chatelier's Principle. (1-4)

Data Tables/Summary: -

Table #1: Signed pre-lab & observations collected during experiment for part 1 / 2

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Table #2: Formatted observations for Part 1

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Table #3: Formatted observations for part 2

Results & Discussion: Logical Explanation (questions answered within): Part 1: When a solution contains chromate ions (CrO^2- 4) the solution is an intense yellow and will turn orange when dichromate ions (Cr2O^2- 7) are present. We see in this in the following reaction: CrO^2- 4 (aq) + 2H+ (aq)  Cr2O^2- 7 (aq) + H2O (l) (reversible). Given that chromate is yellow and dichromate is orange we are told they are in equilibrium in the above reaction. The acidity of the solution will affect the equilibrium by shifting to the left and when there’s a higher level of acidity the equilibrium will shift to the right. In test tube #1 our solution

went from yellow to orange to back to a bright yellow color. When adding drops of HCl to the chromate solution this tells that equilibrium as shift to the right when increasing the hydrogen ion concentration. But when we added drops of NaOH this showed favored toward the reactants side and shifted to left which is why it went back to the color yellow. For test tube #2 our solutions went from faint orange then remained the same then to the color yellow. When we added drops of dichromate and HCl there was no change of the faint orange color because since we added H+ the reaction consumed it and produce more products. Thus, when added our drops of NaOH it helped remove H+ turning the solution yellow and produced more reactants (shifting to the left). Next for test tube #3 our solution went from yellow to no change then to orange. Since no color changed occurred between the adding of chromate and NaOH it tells us that equilibrium shift to the left. Once we added HCl it removed H+ and converted the solution to an orange color which then shifted equilibrium to the right. For test tube #4 our color changes went from starting at orange to neon yellow and back to orange. Since the system lost reactants because NaOH removed H+ it tries to produce more reactants by shifting equilibrium to the left. Once adding HCl it turned the solution to orange shifting equilibrium to the right. For test tube 5 to 8 we started adding Ba2+ onto our solution in which we started noticing precipitation (cloudiness) start occurring except for test tube 8. For this precipitation reaction when then get the balanced equation of: Ba2+ (aq) + CrO^2-4 (aq) + 2H+ ⇌ Cr2O7^2- + H2O (l). Next for test tube #5 my reactions went for a cloudy yellow to a faint orange with no cloudiness. This occurred because we added NaOH before Ba2+ which shifted the equilibrium to the left and produced more chromate leading to a cloudy yellow color. Once we added HCl to the mix it increased H+ which shift equilibrium to the right and concentration of chromate decreased leading to not precipitation. For test tube #6 changes occurred from no changes until we add

NaOH which turned a mixture to a cloudy yellow. This occurred because we added NaOH removed H+ and shifted the equilibrium to the left and NaOH was added after Ba+2. Next for test tube #7 we got a cloudy neon orange color (precipitated) since we know chromate and Ba2+ react with each other. Finally, for test tube #8 we got no precipitation because most of the Ba2+ did not react with dichromate. Part 2: For the second part of the experiment we are shown the reaction of the second equilibrium of: [CoCl4]2- (alc) + 6H2O (l) ⇌ [Co(H2O)6]2+ (alc) + 4Cl- (alc) (reversible), where [CoCl4]2gives a blue color while [Co(H2O)6]2+ gives off a rose-pink color. We also recall that when stresses are placed on systems in equilibrium ca be monitored by observing changes in color and their measured absorbance’s which will give maximum sensitivity for concentration differences. For test tube #2 we observed a light blue color because HCl increased the amount of chlorine anions making the equilibrium shift the left to consume any additional anions. For test tube #3 we obtained a light pinkish color once we added drops of distilled water because water removed chlorine anions, shifting the equilibrium to the right to produce more anions for those lost. Another note is when dropping distilled water to the solution there was a division between to mixtures where the color blue was at top and the pinkish color remained at the bottom of beaker before mixing. For test tube #4 we obtained a precipitated purple color because the chlorine anions reacted with Pb+2 ions (Pb2+ (alc) + 2Cl- (alc)  PbCl2 (s)). For test tube #5 we obtained a dark blue color after placing it into a hot water bath which increased the temperature of the mixture and gave it a dark blue color. Thus, equilibrium shifted to the left. Finally, for test tube #6 we observed a light pink color when we placed the mixture into an ice water bath which decreased the temperature and gave it, its color. Then indicating that the equilibrium shifted to

the right. Going based off what we observed and collected data of, the forward reaction in the equilibrium in part 2 is exothermic. I come to this conclusion because in test tube #5 when we increased the temperature of the mixture the equilibrium shifted to the left. Whereas, test tube #6 when we placed mixture in an ice cold bath it decreased the temperature and shifted equilibrium to the left which goes along with Le Chatelier’s principle. Conclusion: By observing two equilibrium systems and how they shift when a stress like concentration and temperature changes done in parts 1 and 2, we were able to understand Le Chatelier’s Principle and decide whether the equilibrium in part 2 is exothermic or endothermic. References: -

Handout given to us by lecture professor / at the end of class syllabus

Anderson, V. K.; Mathias, E. V. Equilibrium and Le Chatelier's Principle. (1-5)...